Chapter 3 Flashcards
The Bohr Rutherford Model of the atoms is the
Planetary model
Rutherford Contribution
- electrons move in space RANDOMLY around the nucleus
- the electrons account for the volume of the atom, and the nucleus accounts for the mass of an atom (neutrons + protons= atomic mass)
–> Rutherford didn’t know about the neutrons in the nucleus
–> It was his student, Chadwick, who later discovered the neutrons in the nucleus
Bohr Contribution
- Electrons can exist only in a series of “allowed” energy levels or shells of “fixed” energies
- Electrons have to be in a shell, it cannot be “in-between” shells
- Each allowed energy state is given a quantum #: n=1,2,3…
- Lower shells have lower energies. Higher shells have higher energies
= planetary model
If the electron is found in the lowest possible energy level, it is in its ___________________
ground state
To jump from a lower level to a higher level, an electron must ____________________. The electron is said to be in an ________________.
absorb energy, excited state
To fall back to a lower level from a higher level, an electron will …
release/emit the same amount of energy it absorbed. This energy may be observed as light
Which absorption would require more energy?
From level 1-2 OR level 4-5?
From level 1-2
- Rmr the distance between shells decreases as # of shells get greater.
Which absorption would release less energy?
From level 3-1 OR level 6-4?
level 6-4
Electromagnetic Spectrum
- consists of light waves of diff frequencies & wavelengths
Wavelength
λ - lamda
- the distance between any 2 corresponding points on adjacent waves –> crest-to-crest OR trough-to-trough
Frequency
ν- nu
- the # of waves that pass in a specific amount of time at a certain point (per second)
- per second= 1/s= s^-1=hertz
Wavelength & frequency are _______________________ to each other
inversely related
- As wavelength increases, frequency decreases
- As wavelength decreases, frequency increases
electromagnetive spectrum from longest to shortest
Radio waves, microwaves, infrared, visible light, ultraviolet, x-rays, gamma
- visible light is obviously safe (i think)
- the first 3 are safe for use to be around
- the last 3 are not safe for us
- x-rays & gamma r not naturally found on earth
Do shorter wavelengths have higher or lower energy?
Shorter wavelengths= higher energy
longer wavelengths= lower energy
The visible spectrum is
the band of light waves (380-750) that the human can detect.
- 1m = 1,000,000,000 (1x10^9)nm
- colours of visible spectrum from smallest to largest wavelength= V,I,B,G,Y,O,R
- colours of visible spectrum from least to most energetic energy= R,O,Y,G,B,I,V
A quantum is
fixed amount/value of energy
plural=quanta
- energy is quantized or restricted to specific levels –> think of ladder
Emission Spectrum
WRITTEN ONE: an excited e- falls from a higher level to a lower one, releasing fixed amount of energy
TYPED ONE: it is the spectrum of frequency of EM (google: electromagnetic) radiation due to an e- making a transition from higher energy state to a lower energy state
- each element in the periodic table has its own “ladder of energies” or “emission line spectrum” signature
How can we explain the lines observed for excited hydrogen
- when an e- gains a fixed amount of energy, it jumps from a higher level & becomes excited
- when this e- falls back to a lower level, a fixed amount of energy is released as light
Hydrogen emission series
LYMAN: n* –> n=1, UV, not visible to naked eye
BALMER: n* –> n=2, visible, visible to naked eye
PASCHEN: n* –> n=3, IR, not visible to naked eye
BRAKET: n* –> n=4, IR, not visible to naked eye
- RMR ladder analogy with its gaps.
- smaller gap in each series= greater relative wavelengh
–> shorter wavelength=greater energy, greater frequency
Orbits VS Orbitals
ORBITS - e- travel around the nucleus in circular 2D pathway. The distances from the nucleus are fixed. No evidence exists for this model & Bohr’s atomic model of e- is abandoned
ORBITALS- the shape of an orbital is defined by the motion of the e- in that orbital. An orbital does not tell us the exact location of an e- where it is most likely to be found, only the 3D space where a 95% probability of finding an e- in it. Each orbital can hold a max of 2e-. (UNCERTAINTY PRINCIPLE) For diff allowed energy states, diff # & types of orbitals exist which are described by a set of quantum #
The principal quantum # - n
- the energy levels
- shows how far an e- is from the nucleus (relative size of the nucleus
- the higher the #, the greater the energy of the e-
Angular Momentum Quantum # - l
uses a # to represent the types & shapes of orbitals within each quantum level
l= 0 to n-1
l=0
- an “s” spherical orbital
- the cloud’s density is not uniform throughout but is greater near the nucleus & decreases as we move away
- only one “s” orbital in each shell & can hold 2e-
l=1
- 3 “p” perpendicular orbitals –> px, py, pz
- dumbbell shape= the lobes of each p orbital disappear at the origin where the nucleus is located –> meaning possibility of finding an e- at the nucleus is zero
- 3 “p” orbitals are = equal in energy (degenerate) & can hold max of 6 e-
l=2
- 5 “d” diffuse orbitals
- 5 diff degenerate orbitals in diff direction
- can hold a max of 10e-
l=3
- 7 “f” fundamental orbitals
- 7 diff degenerate orbitals in diff direction
- can hold a max of 14e-
The Magnetic Quantum # -ml
- indicates the 3D orientation of an e- (which plane the e- is found in)
-l <= ml<= l
The Magnetic Spin Quantum # - ms
- the direction of e- spin
- each orbital can hold a max of 2 e- with opposite spin ( the Pauli Exclusion Principle)
e- configurations using the condensed format
An atom’s full e- shells are represent by the symbol of the nearest preceding Noble Gas, in square brackets, followed by the e- configurations for the remaining e-
Aufbau Principle
e- will fill lowest energy levels before higher ones are filled
Pauli Exclusion Principle
each orbital holds a maximum of two e- with opposite spin
Hund’s Rule
when there are orbitals of the same energy within a sublevel, electrons will half-full each orbital before pairing up –> this is cuz e- will repel each other and will want to be as far as possible
Electron configuration for ions
cations (+) = e- are lost from the largest (outermost) shell –> highest n value, if n values are the same, then e- are removed from higher energy orbital first
anions (-)= e- are added to the nearest orbitals that are not yet full
Paramagnetic VS diamagnetic
types of magnetism
PARAMAGNETIC
- unpaired electrons
- attracted to a magnetic field
DIAMAGNETIC
- paired electrons
- repelled by a magnetic field
Identifying if an electron configuration is in its ground or excited state
- If valence e- are found in the order that we would predict, then the e- are in their ground state & as close to the nucleus as possible
- if the e- configuration is “out of order” (like if there is a gap), it means that e- are not in their ground state. –> they are in energy levels further from the nucleus than expected, they are in an excited state
Exceptions to the electron configurations (Anomalous cases)
- some elements that do not behave according to their predicted e- configurations
- it turns out that when the “d” orbitals are half full or completely full, the e- can drop closer to the nucleus. –> This means that it is lower energy, or more stable, when the “d” orbitals are half or completely full
- these elements are: Nb, Mo, Cr, Ru, Rh, Pd, Ag, Pt, Au, Cu
- S^2 –> S^1
–> only exception is Pd –> [Kr] 5s^2 4d^8 changes to [Kr]4d^10
